During the latter half of the eighteenth century, most parts of Western Europe were healthily in the midst of the Industrial Revolution and the chemical industry was at the heart of this expansive advance. For more than a century, one of the most demanded class of chemicals had been the alkali carbonates because they were central to the glass, soap, and masonry industries. Potassium carbonate was extracted from the ashes of inland plants and sodium cabonate was similarly produced form the ashes of certain coastal plants. This alkali industry along with growing demand for charcoal and lumber caused massive deforestation in Europe by the turn of the century. Therefore, European industries were increasingly forced to rely on importing sodium and potassium carbonate from trona and related mines in the United States.
The high cost of alkali carbonates soon became so much of a problem that the French Académie des sciences offered a large reward to anyone who could cheaply convert sea salt to sodium carbonate. In 1791, Nicolas LeBlanc succeded in doing just that by first making Glauber's salt from table salt and sulfuric acid.
H2SO4 + 2NaCl --> Na2SO4 + 2HCl
This reaction actually occurs in two steps but the latter only really gets going with a fused mass at high temperatures.
NaHSO4 + NaCl --> Na2SO4 + HCl
Next, the sodium sulfate is reacted with calcium carbonate from limestone and charcoal in a furnace to form sodium carbonate, calcium sulfide, and carbon dioxide.
Na2SO4 + CaCO3 + 2C --> Na2CO3 + CaS + 2CO2
The cooled mass is then lixivated with water and filtered from the insoluble sulfide, charcoal, and calcium carbonate.
LeBlanc's process satisfied the requirements of the Académie des sciences, but sadly the revolutionary government at the time took away his factory and refused to grant him his prize. He would spend the last 15 years of his life on the brink of bankruptcy before shooting himself in 1806.
Today, the LeBlanc process is outdated and has been entirely replaced by the Solvay process with Hao's modification. The key reaction in this brilliantly elegant process is the precipitation of sodium bicarbonate in a concentrated solution of sodium chloride and ammonium bicarbonate.
NaCl + NH4HCO3 --> NaHCO3(s) + NH4Cl
However, the potassium analogy of the Solvay reaction does not work because potassium bicarbonate is too soluble to precipitate out. Therefore, potassium carbonate is now produced by electrolyzing a solution of potassium chloride and subsequently reacting the formed potassium hydroxide with carbon dioxide.
I decided to produce my potassium carbonate through a modified version of the LeBlanc process. I used potassium bisulfate instead of potassium sulfate because this was what was easily left over from my nitric acid distillation.
H2SO4 + KNO3 --> KHSO4 + HNO3
My hypothesis was that it would be fine to use the bisulfate instead of the sulfate in the LeBlanc process except that you would need twice as much of the salt. The proposed reaction looks something like this:
2KHSO4 + 2CaCO3 + 4C --> K2CO3 + 2CaS + 5CO2 + H2O
Also, I planned to run this reaction in a barbeque which probably does not even get to cherry red heat. This was bad news because most sources state that this process was run at temperatures around 1000C, but I was hoping that the reaction would still occur at lower temperatures, albeit less effectively, as long as the carbon could be oxidized.
Potassium carbonate is a chemical that I would like to have relatively large amounts of because it is instrumental in making all other potassium salts. Potassium salts are generally less hygroscopic than their sodium relatives and are more soluble in organic liquids. I have ready access to the chloride and nitrate, but I really need the carbonate for more flexibility in synthesis.
I plan to use a good deal of any potassium carbonate produced to make potassium hydroxide, or potash lye, by reacting it with slaked lime.
K2CO3 + Ca(OH)2 --> CaCO3(s) + 2KOH
I made a couple of key mistakes over the course of this experiment that will adversely affect yields, but at least I proved that the reaction occurs, although perhaps not completely. I first made calcium carbonate by dissolving 22.2g CaCl2 and 33.6g NaHCO3. The following reaction occurred as carbon dioxide was given off.
CaCl2 + 2NaHCO3 --> CaCO3(s) + 2NaCl + CO2 + H2O
I heated the mixture until boiling in order to digest the calcium carbonate precipitate because it was a very fine white powder. Although I added enough to theoretically make 20g, I only ended up with 13g of calcium carbonate after thoroughly drying it. I added the 13g to 20g of KHSO4 so I ended up with too little calcium carbonate in the mixture. However, the KHSO4 probably had some water of crystallization that came along with it and some potassium nitrate as well. In hindsight, I should have dried this mixture of salts in the oven to get a more certain idea of what I was working with.
At the beginning of heating in a tuna fish can, a relatively small amount of yellow-orange nitrogen dioxide fumes were given off due to the reaction of potassium bisulfate and left over potassium nitrate. This reaction effectively got rid of any nitrate impurity and some of the bisulfate was converted to the sulfate. Over time, the mixture slowly, and at first unnoticeably, turned from an initial light grey to more of a creamy white.
After 100 minutes of heating, I tested a picayune bit of my product with a drop of sulfuric acid and I got the characteristic smell of hydrogen sulfide gas from the calcium sulfide that formed. I then thoroughly mixed the sludge with about 450mL of water and filtered the insoluble material off. I then tested a bit of the solution with another drop of sulfuric acid and this fizzed but did not smell of eggs, letting me know that there was at least some potassium carbonate in there.
At this point, I suspect that I had a mixture of potassium sulfate and potassium carbonate. I had previously witnessed some minor bubbling occur when I poured water of the sludge for the first time. Potassium bisulfate left over from the reaction reacted with the formed potassium carbonate to form carbon dioxde and potassium sulfate. If there was stoichiometrically more potassium bisulfate in there then the carbonate, then that could have completed killed my yields.
2KHSO4 + K2CO3 --> 2K2SO4 + CO2 + H2O
The carbon dioxide formed also means that some of the carbonate was converted into the bicarbonate which did not help my fractional crystallization attempts.
I began fractional crystallizion of the solution of potassium sulfate and potassium carbonate with perhaps a bit of potassium bicarbonate by boiling it down to 150mL. This gave some precipitate even at boiling temperature because only 1g of potassium sulfate dissolves in 8.4mL of water at room temperature. I performed the sulfuric acid test again, this time with the decanted precipitate, and it tested positive for carbonate/bicarbonate ion so I knew that I had already crystallized out more than I needed to. I then continued to boil the now 130mL solution down and nothing precipitated out until it was at about the 70mL level. This was a good sign because potassium carbonate is very soluble at 25C (111g/100mL). Probably, potassium bicarbonate had already precipitated out along with the sulfate because its solubility is 36g/100mL at 25C. When the solution reached 60mL, I stuck it in the freezer and then filtered it out. I repeated this process when the water level was at about 30mL and 10mL. With this multiple crystallization method, I prevented the solution from splattering during boiling and was able to assay my product each time.
Next, I put the wet yield in the oven to dry it and I ended up with 3.7g of K2CO3 with a slight yellow tint.
Overall, I am fairly happy that at least I can replicate the LeBlanc process in this manner. I think that if I use a larger quantity of reagents and if I make sure that potassium bisulfate is not in excess, I can achieve a better percentage yield. A 37% yield is definitely nothing to be proud off.